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Why does a reaction push back when you change it?

The GCSE picture: reversible reactions, dynamic equilibrium and Le Chatelier’s principle.

  1. A sealed tube of brown gas10 cm
  2. A pocket of gas1 µm
  3. A few hundred molecules50 nm
  4. Two molecules joining1 nm
2NO2(g) ⇌ N2O4(g) ΔH = −57 kJ mol−1
What the simulation shows, in words

Brown dots are molecules of nitrogen dioxide, NO₂. Pale joined pairs are dinitrogen tetroxide, N₂O₄, which is colourless. When two NO₂ meet they may join; any N₂O₄ may split back into two NO₂. After a few seconds the numbers settle down, with both reactions still going on: that is dynamic equilibrium. Heating makes more N₂O₄ split, so the gas turns browner. Shrinking the box squeezes the molecules together and more pairs form, so after a first darkening the colour fades. Adding NO₂ makes more pairs form too. The graph tracks both numbers over time.

Two reactions at once

Nitrogen dioxide is a brown gas. Two of its molecules can join to make one molecule of dinitrogen tetroxide, which is colourless. That molecule can also split back into two:

2NO2(g) ⇌ N2O4(g)

In a sealed tube both reactions go on all the time. At first, with lots of NO₂, pairs form quickly. As N₂O₄ builds up, more of it splits. Soon the two reactions run at the same rate. From then on the amounts stop changing, even though molecules are still joining and splitting every instant. That is a dynamic equilibrium. It only happens in a closed system, where nothing can get in or out.

“Equilibrium” doesn’t mean equal amounts. It means the forward and reverse rates are equal, so the amounts are constant.

Pushing it, and how it pushes back

Le Chatelier’s principle says that if you change the conditions of a system at equilibrium, the position of equilibrium moves to counteract the change. Each of these shows up in the colour of the tube.

Temperature. Joining two NO₂ molecules gives out heat: the forward reaction is exothermic. So the reverse, splitting, takes heat in. Heat the tube and the equilibrium moves in the endothermic direction, to the left, to use up some of the extra heat. More NO₂ means a browner gas. Cool it and the gas turns paler.

Pressure. The left side has two molecules of gas, the right side has one. Squash the gas into a smaller space and the pressure goes up. The equilibrium moves towards the side with fewer gas molecules, to the right, which lowers the pressure again. In the simulation, shrink the box: the gas first looks darker (the same molecules in less space), then fades a little as pairs form.

Concentration. Add more NO₂ and some of it joins up, so the equilibrium moves to the right to use up part of what you added.

A catalyst speeds up both reactions equally. Equilibrium arrives sooner, but its position doesn’t move.

Why it matters

Industry uses this all the time. Ammonia for fertilisers is made by a reversible reaction, N₂ + 3H₂ ⇌ 2NH₃, and the conditions are chosen by thinking through exactly these effects, then balancing the yield against the rate and the cost.

Safety

Nitrogen dioxide is toxic. Sealed tubes of it are only used in school demonstrations by teachers. Never try to make or release it.